View source for Ferrate(VI)

{{merge to|Potassium ferrate|discuss=Talk:Potassium ferrate#Merge proposal|date=April 2025}}
{{Short description|Ion}}
{{Chembox
| verifiedrevid = 443819490
| ImageFile = Ferrate ion.svg
| ImageFile_Ref = {{chemboximage|correct|??}}
| ImageSize = 121
| ImageName = Aromatic skeletal formula of ferrate
| ImageFile1 = Ferrate and permanganate solution.jpg
| ImageCaption1 = Solutions of ferrate (left)<BR> and [[permanganate]] (right) 
| IUPACName = Ferrate(VI)
| SystematicName = Tetraoxoironbis(olate){{citation needed|date = March 2011}}
| OtherNames = [FeO<sub>4</sub>]<sup>2-</sup>
|Section1={{Chembox Identifiers
| PubChem = 25000034
| ChemSpiderID = 21865127
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 30992
| Gmelin = 2055
| SMILES = [O-][Fe]([O-])(=O)=O
| SMILES1 = [O-][Fe](=O)(=O)[O-]
| StdInChI = 1S/Fe.4O/q;;;2*-1
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| InChI = 1/Fe.4O/q;;;2*-1/rFeO4/c2-1(3,4)5/q-2
| StdInChIKey = XGBDPAYTQGQHEW-UHFFFAOYSA-N
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| InChIKey = XGBDPAYTQGQHEW-WTZHFVRHAI
}}
|Section2={{Chembox Properties
| Formula = [FeO<sub>4</sub>]<sup>2-</sup>
| MolarMass = 119.843 g mol<sup>−1</sup>
}}
}}

'''Ferrate(VI)''' is the [[inorganic chemistry|inorganic]] [[anion]] with the [[chemical formula]] [FeO<sub>4</sub>]<sup>2−</sup>. It is [[Photosensitivity#Interpretation in chemistry|photosensitive]], contributes a pale violet colour to compounds and solutions containing it and is one of the strongest water-stable oxidizing species known. Although it is classified as a [[basicity|weak base]], concentrated solutions containing ferrate(VI) are corrosive and attack the skin and are only stable at high [[pH]]. It is similar to the somewhat more stable [[permanganate]].

== Nomenclature ==

<!--The systematic name ''tetraoxidoferrate(2−)'', a valid [[IUPAC]] name, is constructed according to the additive nomenclature. In the case where the [[multiplicity (chemistry)|multiplicity]] is required to be specified, a more precise name is used. Singlet ferrate(VI)'s additive name remains unchanged. Triplet ferrate(VI), which is the ground state, is named ''tetraoxidoferrate(2•2−)''. -->
{{See also|Ferrate}}

The term ''ferrate'' is normally used to mean ferrate(VI), although it can refer to other [[iron]]-containing [[anion]]s, many of which are more commonly encountered than salts of [FeO<sub>4</sub>]<sup>2−</sup>.  These include the highly reduced species [[disodium tetracarbonylferrate]] {{chem2|Na2[Fe(CO)4]}}, {{chem2|K2[Fe(CO)4]}} and salts of the iron(III) complex tetrachloroferrate [FeCl<sub>4</sub>]<sup>−</sup> in [[1-Butyl-3-methylimidazolium tetrachloroferrate]].  Although rarely studied, ferrate(V) [FeO<sub>4</sub>]<sup>3−</sup> and ferrate(IV) [FeO<sub>4</sub>]<sup>4−</sup> [[oxyanion]]s of iron also exist.  These too are called ferrates.<ref name="hill">{{cite book
| title = Chemistry in context
| author1 = Graham Hill
| author2 = John Holman
| edition = 5th
| publisher = Nelson Thornes
| year = 2000
| isbn = 0-17-448276-0
| page = 202
}}</ref>

==Synthesis==

Ferrate(VI) salts are formed by [[redox|oxidizing]] iron in an [[aqueous]] medium with strong oxidizing agents under [[base (chemistry)|alkaline conditions]], or in the solid state by heating a mixture of iron filings and powdered potassium nitrate.<ref name="sharma">{{cite book
| title = Text Book Of Coordination Chemistry
| author = R. K. Sharma
| publisher = Discovery Publishing House
| year = 2007
| isbn = 978-81-8356-223-2
| pages = 124–125
}}</ref>

For example, ferrates are produced by heating [[iron(III) hydroxide]] with [[sodium hypochlorite]] in alkaline [[Solution (chemistry)|solution]]:<ref name="wulfsberg"/>

:2 {{chem|Fe(OH)|3}} + 3 {{chem|OCl|-}} + 4 {{OH-}} → 2 {{chem|[FeO|4|]|2-}} + 5 {{H2O}} + 3 {{chem|Cl|-}}

The anion is typically precipitated as the [[barium|barium(II)]] salt, forming [[barium ferrate]].<ref name="wulfsberg"/>

==Properties==

Fe(VI) is a strong oxidizing agent over the entire pH range, with a reduction potential (Fe(VI)/Fe(III) couple) varying from +2.2&nbsp;V to +0.7&nbsp;V versus [[Standard hydrogen electrode|SHE]] in acidic and basic media respectively.

:{{chem|[FeO|4|]|2-}} + 8 {{H+}} + 3 e<sup>−</sup> {{eqm}} {{chem|Fe|3+}} + 4 {{H2O|nolink=y}}; ''E''<sub>0</sub> = +2.20 V (acidic medium)
:{{chem|[FeO|4|]|2-}} + 4 {{H2O|nolink=y}} + 3 e<sup>−</sup> {{eqm}} {{chem|Fe|(OH)|3}} + 5 {{Chem|OH|-}}; ''E''<sub>0</sub> = +0.72 V (basic medium)

Because of this, the ferrate(VI) anion is unstable at neutral<ref name="sharma" /> or acidic [[pH]] values, decomposing to iron(III):<ref name="wulfsberg">{{cite book
| title = Principles of descriptive inorganic chemistry
| author = Gary Wulfsberg
| publisher = University Science Books
| year = 1991
| isbn = 0-935702-66-0
| pages = 142–143
}}</ref> The reduction goes through intermediate species in which iron has oxidation states +5 and +4.<ref name="wiberg"/> These anions are even more reactive than ferrate(VI).<ref name="bergeron"/> In alkaline conditions ferrates are more stable, lasting for about 8 to 9 hours at pH 8 or 9.<ref name="bergeron">{{cite book
| title = The Development of Iron Chelators for Clinical Use
| editor = Raymond J. Bergeron
| author = Gary M. Brittenham
| publisher = CRC Press
| year = 1994
| isbn = 0-8493-8679-9
| pages = 37–38
}}</ref>

Aqueous solutions of ferrates are pink when dilute, and deep red or purple at higher concentrations.<ref name="wiberg">{{cite book
| title = Inorganic chemistry
| author1 = Egon Wiberg
| author2 = Nils Wiberg
| author3 = Arnold Frederick Holleman
| publisher = Academic Press
| year = 2001
| isbn = 0-12-352651-5
| pages = 1457–1458
}}</ref><ref>{{cite book
| title = Oxford dictionary of chemistry
| editor = John Daintith
| edition = 5th
| publisher = Oxford University Press
| year = 2004
| isbn = 0-19-860918-3
| page = 235
}}</ref> The ferrate ion is a stronger [[oxidizing agent]] than [[permanganate]],<ref name="mackay">{{cite book
| title = Introduction to modern inorganic chemistry
| author1 = Kenneth Malcolm Mackay
| author2 = Rosemary Ann Mackay
| author3 = W. Henderson
| edition = 6th
| publisher = CRC Press
| year = 2002
| isbn = 0-7487-6420-8
| pages = 334–335
}}</ref> and oxidizes [[ammonia]] to molecular [[nitrogen]].<ref>{{cite web | url = https://kb.osu.edu/dspace/bitstream/1811/36335/1/OH_WRC_444.pdf | title = Oxidation of Ammonia in Water by Ferrates(VI) and (IV) | author = Karlis Svanks | date = June 1976 | page = 3 | publisher = Water Resources Center, Ohio State University | access-date = 2010-05-04 }}</ref>

The ferrate(VI) ion has two [[unpaired electron]]s and is thus [[paramagnetism|paramagnetic]]. It has a [[tetrahedral molecular geometry]], isostructural with the chromate and permanganate ions.<ref name="wiberg" />

== Applications ==
Ferrates are excellent [[disinfectant]]s, and are capable of removing and destroying [[virus]]es.<ref>{{cite book
| title = Environmental chemistry
| author = Stanley E. Manahan
| edition = 8th
| publisher = CRC Press
| year = 2005
| isbn = 1-56670-633-5
| page = 234
}}</ref> They are also of interest as potential as an environmentally friendly water treatment chemical, as the byproduct of ferrate oxidation is the relatively benign iron(III).<ref>{{Cite journal|last1=Sharma|first1=Virender K.|last2=Zboril|first2=Radek|last3=Varma|first3=Rajender S.|date=2015|title=Ferrates: Greener Oxidants with Multimodal Action in Water Treatment Technologies|url=https://pubs.acs.org/doi/10.1021/ar5004219|journal=Accounts of Chemical Research|language=en|volume=48|issue=2|pages=182–191|doi=10.1021/ar5004219|pmid=25668700|issn=0001-4842|url-access=subscription}}</ref>

Sodium ferrate ({{chem2|Na2FeO4}}) is a useful reagent with good selectivity and is stable in aqueous solution of high pH, remaining soluble in an aqueous solution saturated with sodium hydroxide.{{Citation needed|date=October 2020}}

==See also==
* [[High-valent iron]]
* [[Potassium ferrate]]
* [[Barium ferrate]]

==References==
{{reflist}}
{{Iron compounds}}

[[Category:Ferrates| ]]
[[Category:Oxidizing agents]]
[[Category:Transition metal oxyanions]]
[[Category:Oxometallates]]